Chemistry · Chemistry of the elements · Grade 7

Oxygen and ozone

Why does a glowing splint burst into flame in oxygen, and what does the ozone layer protect us from?

The situation

In a hospital, a patient who has difficulty breathing is given pure oxygen through a mask — it comes from a sky-blue steel cylinder. Where does this oxygen come from? In a school laboratory it can be obtained by heating potassium permanganate and collecting it over water in a test tube: if a glowing splint is lowered into this test tube, it bursts into flame. Why does an iron wire not burn in air, but burn in oxygen, throwing out sparks? And the other form of oxygen — ozone — which of the Sun's rays does it protect us from at an altitude of 20–30 km?

Topic

Oxygen and ozone. Oxygen cylinder in a hospital and the ozone layer: heating permanganate to obtain oxygen — a glowing splint relights

Key ideas

Preparing and collecting oxygen
Glowing splint test
Ozone layer

Formulas

2 KMnO_{4} \overset{\overset{}{t}}{\to} K_{2} MnO_{4} + MnO_{2} + O_{2} ↑

In the laboratory: heating potassium permanganate — it decomposes from ≈ 240 °C.

2 H_{2} O_{2} \overset{\overset{}{MnO_{2}}}{\to} 2 H_{2} O + O_{2} ↑

Hydrogen peroxide decomposes at room temperature; MnO₂ is a catalyst and is not itself used up.

n (O_{2}) = \frac{1}{2} n (KMnO_{4}), V = n V_{m}

Volume of oxygen released: at STP $V_{m}$ = 22.4 L/mol, at 20 °C ≈ 24.1 L/mol.

3 Fe + 2 O_{2} \overset{}{\to} Fe_{3} O_{4}

Iron burns in oxygen throwing out sparks — iron scale forms.

I = I_{0} e^{- σ N / \sin h}

UV light that has passed through the ozone layer: $σ$ — absorption cross-section of ozone, $N$ — amount of ozone, $h$ — height of the Sun.

Theory

Oxygen in nature and its physical properties

Oxygen (O) is the most abundant element in the Earth's crust: nearly half of the mass of the crust. 20.95 % of the volume of air is the simple substance O₂, and 89 % of the mass of water is oxygen. O₂ is a colourless, odourless gas, $M_{r}$ = 32; slightly heavier than air: $D_{a i r} = 32 / 29 \approx 1.1$ , density at STP 1.43 g/L. It is sparingly soluble in water: in water saturated with air at 20 °C ≈ 9 mg/L — fish breathe this oxygen. It liquefies at −183 °C; liquid oxygen is pale blue and is attracted by a magnet.

Preparation: in the laboratory and in industry

In the laboratory oxygen is obtained by decomposing substances rich in oxygen. Potassium permanganate decomposes when heated: $2 KMnO_{4} \overset{}{\to} K_{2} MnO_{4} + MnO_{2} + O_{2} ↑$ . If a little manganese(IV) oxide is added to a hydrogen peroxide solution, it froths at room temperature: $2 H_{2} O_{2} \overset{}{\to} 2 H_{2} O + O_{2} ↑$ . MnO₂ speeds up the reaction but is not itself used up — it is a catalyst. Berthollet's salt (potassium chlorate) also gives oxygen when heated with MnO₂: $2 KClO_{3} \overset{}{\to} 2 KCl + 3 O_{2} ↑$ . In industry oxygen is obtained from air: air is cooled and liquefied, then slowly evaporated — nitrogen (−196 °C) boils first, oxygen (−183 °C) later. Oxygen is stored in steel cylinders at ≈ 15 MPa pressure: a 40-litre cylinder holds ≈ 6 m³ of gas. Uses: in medicine for breathing, in metallurgy for smelting cast iron and steel (air is enriched with oxygen), for welding and cutting metals (the acetylene–oxygen flame), and liquid oxygen as the oxidiser for rocket fuel.

Practical: collecting and identifying oxygen

O₂ is sparingly soluble in water, so it can be collected by displacement of water: a test tube filled with water is inverted in a trough, and the delivery tube is brought under it. O₂ is heavier than air, so in the displacement of air method the test tube is held mouth upwards. Oxygen is recognised by a glowing splint: in a gas rich in oxygen it bursts into flame. The first bubbles are the air in the apparatus (the heated air expands and comes out), so the gas in the first test tube is not pure. Calculation: $n (O_{2}) = \frac{1}{2} n (KMnO_{4})$ , $V = n V_{m}$ : from 3 g of KMnO₄ theoretically 9.49 mmol, that is 213 mL (STP) of oxygen is released.

Chemical properties: combustion and oxidation

Oxygen is a strong oxidising agent: many simple substances combine with it to form oxides. Carbon: $C + O_{2} \overset{}{\to} CO_{2}$ ; sulfur (blue flame): $S + O_{2} \overset{}{\to} SO_{2}$ ; phosphorus (white smoke): $4 P + 5 O_{2} \overset{}{\to} 2 P_{2} O_{5}$ ; magnesium (white light): $2 Mg + O_{2} \overset{}{\to} 2 MgO$ ; iron (sparks): $3 Fe + 2 O_{2} \overset{}{\to} Fe_{3} O_{4}$ . Burning in pure oxygen is much more violent than in air: in air only one fifth of the molecules are O₂, whereas in pure oxygen almost all are — oxygen molecules hit the surface of the fuel many times more often, and heat is released faster. That is why an iron wire does not burn in air but burns in oxygen. The combination of a substance with oxygen is oxidation; rusting and respiration are slow oxidation.

Ozone and the ozone layer

Ozone O₃ is an allotropic form of oxygen: one element forms two different simple substances. Ozone is a pale blue gas with a sharp smell, a much stronger oxidising agent than O₂. It forms in an electric discharge (lightning, an ozoniser): $3 O_{2} \overset{}{\to} 2 O_{3}$ ; in the stratosphere, the Sun's short-wavelength (shorter than 242 nm) UV light splits O₂ into atoms, and an atom combines with O₂ to give O₃. Most of the ozone is at an altitude of 15–35 km — the ozone layer. It absorbs light of wavelengths shorter than 320 nm: UV-C completely and most of UV-B. If all the ozone were collected at STP, the layer would be only ≈ 3 mm thick (300 Dobson units). The chlorine in freons (CFCs) destroys ozone: where there is less than 220 DU it is an ozone hole; the 1987 Montreal Protocol phased out freons. Ozone is used to disinfect drinking water (ozonation); ozone in the air near the ground, however, is harmful to the respiratory tract.

Biological importance, safety and the model

In respiration, oxygen oxidises nutrients and releases energy: $C_{6} H_{12} O_{6} + 6 O_{2} \overset{}{\to} 6 CO_{2} + 6 H_{2} O$ ; plants do the reverse in photosynthesis, returning oxygen to the air — this is the oxygen cycle in nature. In medicine, pure oxygen is given in respiratory failure. Safety: wear safety goggles, take small amounts of the substances; before putting out the spirit lamp, take the delivery tube out of the water — otherwise the cooling gas will suck the water back and the hot test tube will crack; burn sulfur and phosphorus in a fume cupboard, use only red phosphorus, and do not stare at burning magnesium. Model: KMnO₄ decomposes according to the school equation (in reality slightly more O₂ is released); the gas in the apparatus and in the test tube mixes uniformly; a glowing splint ignites when the O₂ fraction exceeds 50 % (a model threshold; the real threshold depends on the splint); burning rates are schematic; ozone: the Bouguer–Lambert–Beer law, cross-sections at 298 K (Molina and Molina, 1986), scattering and clouds are not taken into account; the atmosphere section is at a scale of 1 m = 50 km, objects on the ground are not to scale.

Did you know?

Oxygen was discovered independently by the Swedish chemist Carl Scheele (1771–1772) and the English scientist Joseph Priestley (1774): Priestley heated mercury(II) oxide with a lens. The name “Oxygène” (“acid-former”) was given by the French chemist Antoine Lavoisier.
Liquid oxygen is pale blue and paramagnetic: if it is poured between the poles of a strong magnet, it does not flow away but hangs between the poles. Liquid nitrogen does not do this.

Inside the experiment

The experiment starts from a real-life situation: you predict the result, change the controls, run the experiment and draw a conclusion. Inside are missions, theory, a quiz and a guided tour in which the camera moves and explains the key points.

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Chlorine and hydrochloric acid — Grade 8
The oxygen group and sulfur — Grade 8
Sulfuric acid and the contact process — Grade 8
Nitrogen, ammonia and nitric acid — Grade 8